BORAX TO BORANES

Unstable

Intermediates

in

the

Pyrolysis of Diborane

1

2

3

J. R.MORREY, A. B.JOHNSON, YUAN-CHINFU, and GEORGE RICHARD HILL

BORAX TO BORANES Downloaded from pubs.acs.org by UNIV OF CALIFORNIA SAN DIEGO on 12/13/16. For personal use only.

Fuel Technology Department, University of Utah, Salt Lake City, Utah

Pyrolysis of diborane was carried out in a borosilicate glass cell designed to freeze out products that would condense at —78.5°C. Between 100° and 130°C. pentaborane(11) was the main condensable product. The stoichiometry of such a reaction caused a decrease in total pressure, as pentaborane(11) exhibits little vapor pressure at the condensation temperature: 5 B H -->2 B H + 4 H . Without exception, an initial increase in pressure was observed. This pressure increase was not due to fast initial decomposition of pentaborane(11) to polymer or pentaborane(9), nor to slow attainment of temperature equilibrium in the reaction vessel. An alternative explanation appeared to be a high initial concentration of unstable intermediates in the formation of pentaborane(11). To test this hypothesis, a differential infrared technique was used. 2

6

5

11

2

S e v e r a l studies have been made of the kinetics of diborane pyrolysis (1-4)· I presentation of these studies the investigators have postulated various unstable intermediates on the basis of diborane disappearance—i.e., B H , B H , B H , B H , etc. Several investigators have postulated borane ( B H ) to be i n equilibrium with diborane; however, to the authors' knowledge, positive physical detection has not yet been made. n t

3

3

7

3

Ô

4

n

e

8

3

D u r i n g the general studies of diborane pyrolysis i n this laboratory, a series of experiments using a borosilicate glass cell (Figure 1) was carried out. This cell was designed to allow all condensable products to freeze out i n a small mushroom-type condenser located at the top of the cell. T o accomplish this the cup surrounding the condenser was cooled with d r y ice and acetone. A pressure transducer (Statham Laboratories, M o d e l Ρ 24-2A-500-P10A) was attached to the capillary leading from the cell, allowing the total pressure to be followed continuously as a function of time at constant temperature. I t had previously been found that, using the "freeze-out method," i t was possible to effect large percentage conversion to pentaborane (11). This being possible up to 130°C, i t appeared reasonable to attempt a kinetic study, as freezing out the first condensable product would minimize the complications arising from the formation of polymer or higher hydrides. The total pressure vs. time plot for a typical "condensation" experiment is given 1 8 8

Present address, General Electric Co., Richland, Wash. Present address, Wright-Patterson Development Center, Dayton, Ohio. Present address, Phillips Petroleum Co., Bartlesville, Okla. 157


158

ADVANCES IN CHEMISTRY SERIES

Figure 1. Borosilicate glass reaction cell in Figure 2. K n o w i n g that pentaborane(ll) was formed i n high yield, the authors were not surprised that the pressure decreased with t i m e : 5B H '-> 2B6H + 4 H 2

e

n

2

Formation of tetraborane would likewise decrease the pressure, whereas pentaborane (9) would increase the pressure : 2B H« —> B4H10 -f- H2 2

5B He —> 2

BJHJ

-}-

6H

2

T I M E IN MINUTES

Figure 2.

Pyrolysis of diborane

T o t a l p r e s s u r e vs. t i m e


MORREY ET AL.

159


(In each case, the higher hydrides are i n a condensed state.) A n initial increase i n pressure was not expected; nevertheless i t occurred without exception when conditions were favorable for the formation of pentaborane(ll) as the principal product. Initially this was explained b y assuming that two competing rate processes take place : ( A ) diffusion of pentaborane(ll) to the condensation trap and ( B ) further decomposi tion of pentaborane(11) to polymer or pentaborane(9). I t was supposed that during the first minutes of the run, ( B ) was faster than ( A ) ; later, because of possible back reactions involving hydrogen, ( A ) became important. T o test the hypothesis, two duplicate experiments were carried out—i.e., a l l conditions were duplicated except the time allowed for reaction. One experiment was quenched upon attaining the maximum pressure, whereas the other was allowed to continue until the pressure decrease was well under way. Subsequent to the quenching, the condensable products were fractionated from hydrogen and remaining Table I. I n i t i a l Pressure at 2 5 ° C , Atm. 0.0549 0.0535

Experiment Duration, Sec. 6960 2100

Effect of Time on Pyrolysis of Diborane Diborane Decomposed,

Stoichiometry»

%

W

X

Y

Z

65.5 37.0

1.04 1.03

0.082 0.070

0.270 0.270

0.024 0.021

•—•

Q

8.3 ± 0 . 5 Ind.*

S 0.01 0.02

• B t H e = WHi + X B * H » + Γ Β β Η η + Z B i o % + SB4H10. T o o little p o l y m e r f o r m e d t o determine the B / H r a t i o . b

diborane, then were determined by infrared spectroscopy. Table I summarizes the results obtained. I n the table the polymer is arbitrarily designated as B H Q . Some of the early determinations of molecular weight of polymers made b y Stock indicated an average of ten boron atoms per molecule. T h e temperature of both runs was 411.1°K. The analysis of products of both experiments showed no appreciable difference in stoichiometry. Unless fortuitous, the initial pressure rise could not be attributed to a high initial rate of pentaborane(ll) decomposition. Subsequently, experiments were carried out with a thermocouple suspended i n the reaction cell to see whether equilibrium was reached before the pressure maximum occurred. They showed that the reaction cell reached temperature equiUbrium within 4 to 5 minutes under the most adverse conditions and remained at equiUbrium throughout the remainder of the run. The above conclusions led to the postulate that the initial pressure increase was attributable to the initial formation of unstable intermediates. T o test this hypothesis, the following experiment was conceived. 1 0

Two infrared cells were filled with a pressure of diborane at room temperature such that when the 10-cm. borosilicate glass cell was placed i n the reference beam and the brass, heatable cell (6.59 cm.) i n the sample beam of a n infrared spectrometer (Perkin-Elmer M o d e l 21), the diborane absorbance i n the sample beam exactly offset that of the reference beam. The objection has been raised that the initial pressure increase might have been due to a build-up of pentaborane(ll) i n the gas phase; the rate of production of pentaborane (11) might have been faster than either the rate of diffusion and con densation or the further decomposition of pentaborane (11). T h e authors are well aware of this possibility. However, because at least four competing rate processes are involved, all with intricate interdependence, i t would be difficult to establish this point experimentally. I n any case, the positive detection of unstable intermediates has been made and offers at least one substantial explanation of the initial pressure increase. If the absorbances of the reference and sample beams are simultaneously strong enough, the detecting thermocouple will be rendered insensitive. I n the experiments here reported, the pressure i n the borosilicate glass cell was about 60 m m . of mercury at room temperature and that of the brass cell was about 88 m m . of mercury. T h e response mechanism of the spectrometer was still operative but sluggish on the



160

Figure 3.

Dismantled

brass infrared cell

strongest peak (6.23 microns). However, the fast scanning rate necessary during preliminary investigations rendered the response mechanism essentially inoperative in some wave length regions. This condition can be readily detected by momentarily closing the shutter to one of the beams and observing the pen deflection; it must be guarded against when applying this technique. The brass cell pictured i n Figure 3 consists of two cells, the inner cell being wound with a resistance thermometer for temperature control, and a Nichrome heating element. The outside cell is primarily for thermal insulation. The resistance thermometer is one of the resistance arms of a Wheatstone bridge i n the temperature controller. After compensation had been accomplished, the brass cell was rapidly heated (see Figure 4 ) . I t took about 3 minutes to attain reaction temperature. The

TIME ( M I N U T E S )

Figure 4.

Heating curve for infrared cell

MORREY ET AL.

161


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LENCTH, MICMONt

Figure 5.

Infrared

spectra

spectra subsequently obtained are given i n Figures 5 and 6. They were taken at approximately 157° and 126.2°C, respectively. Spectra A, B, and C (Figure 7), and 8 are of tetraborane, pentaborane(9), and pentaborane ( 11 ) and diborane, respectively. They were included for reference. Discussion of Infrared

Peaks

After the cell had been allowed to heat for 1 to 2 minutes, a scan of the 2- to 12-micron region was made at the rate of 2 microns per minute. It may be noticed (5-3, 4; 6-3, 4, 5) that absorption at 2.8 to 2.85, 3.73 to 4.05, 5.5 to 5.8, 6.0 to 6.75, and 8.0 to 9.0 microns was due to compounds or species found in the heated cell which were not found i n the borosilicate glass cell. A s would be expected, the diborane decomposed very rapidly. Decomposition produced over compensation by the reference beam at 2.7, 3.7 to 4.05, 5.15 to 5.5, 6.05 to 6.5, 8.15 to 9, and a broad peak from 9.3 to 11.2 microns. I n other words, the disappearance of reactant is given by maxima i n the curves whereas the appearance of any products is given by minima.

WAVE

LENGTH, MICRONS

Figure 6.

Infrared

spectra



162

WAVE LENGTH, MICRONS

Figure 7. Α . B4H10

Infrared spectra B5H9 C. BeHn

Β.

The peak at 2.82 microns has no analog i n the known volatile hydrides. I t might be argued that i t is simply an overshoot from the maximum caused b y diborane a b sorbing at 2.7 microns. However, spectrum 5-4 verifies that no overshoot takes place.


MORREY ET AL.

163

Pyroiysis of Diborane

The region between 3.75 and 4.02 microns represents an insensitive zone; the recorder pen was simply drifting as a result of inadequate thermocouple response. However, at the edges of this region (3.70 and 4.06 microns) the pen was sensitive. Diborane, p e n t a b o r a n e ( l l ) , and tetraborane a l l absorb at 4.05 microns. They also absorb i n the region from 3.75 to 3.95 microns. One might suspect therefore that the minimum at 4.05 microns (5-3, 4, and 6-3, 4, 5) is due to pentaborane(ll) a n d / o r tetraborane. However, other regions (4.65, 4.90 microns) give ample evidence that neither pentaborane(ll) nor tetraborane was present i n sufficient quantities to account for the minimum at 4.05 microns. F o r example, the absorbance at 4.05 microns predicted for B H would not exceed 0.02, whereas actual absorbance was 0.135. Absorbance is defined as (D = l o g IQ/I) where I is the intensity of the incident light and / is the intensity of the light leaving the cell. The peak at 5.54 microns probably is due to pentaborane(9). Pentaborane(9) also exhibits a peak at 6.15 microns (£>β.ΐδ/Αί.54 — 0.221). F r o m the absorbances, however, i t is easily seen that the strong 6.05-micron absorbance cannot be attributed to pentaborane(9) exclusively; neither can i t be attributed to pentaborane(ll) or tetraborane. The peak at 6.55 microns is unique, having no counterpart i n any of the common volatile boron hydrides. I t is not due to any polymeric compound. I n spectrum 5-3, another insensitive zone occurs between 6.05 and 6.40 microns. H o w ever, because of the disappearance of most of the diborane i n the sample cell, this same region i n spectrum 5-4 is not an insensitive zone. The region between 8 and 9 microns exhibits a peak having a minimum at a p proximately 8.4 microns. However, the peak due to the greater amount of diborane in the reference cell (overcompensation) has cancelled a l l but the edges of this peak. I f one examines the pure diborane peak (Figure 8) at this point, a strong maximum appears at 8.55 microns with two rather broad heavy peaks at either side at 8.4 and 8.72 microns. Examination of the spectra (5-3, 4 and 6-3, 4) reveals complete disap pearance of the 8.4-micron and a diminishing of the 8.55-micron peak. This can only be explained b y a cancellation due to a peak i n the opposite direction. Spectrum 6-5 shows the reappearance of the 8.4-micron and strengthening of the 8.55-micron peaks of the reference cell. This gives evidence that the peak due to intermediates disappears with time. Examination of the spectrum (5-4) shows a tailing off of the "overcompensa t i o n " peak between 6.0 and 6.3 microns. I t should, however, have approximately equal absorption at 6.15 and 6.4 microns. T h a t i t does not is explained b y the continual existence of the strong minimum at approximately 6.15 microns. I t is 5

n

1 0

0



164

noticed i n addition that the peak at 6.55 microns diminishes with time—see, for example, 5-4 and 6-4, 5. I n view of the fact that the observed absorption peaks (minima) for the unstable intermediates are i n each case close to absorption peaks of diborane, one might ques tion whether pressure broadening of the hot diborane could account for some of the peaks ascribed to unstable intermediates. There is evidence to believe that this effect is negligible; spectrum 5-2 should have shown the effect if not negligible since the pressure i n the sample cell had to be 1.5 times that i n the reference cell for balancing, even at room temperature. Furthermore, as soon as the diborane i n the sample cell was heated, the diborane started to decompose into hydrogen, polymer, and a small amount of pentaborane(9). One would not expect a hydrogen molecule to perturb the rotational spectra of a diborane molecule as much as the diborane molecule it replaced. Predicted Bond Constants for BH Assuming Infrared Frequencies Experimentally Obtained 3

Α ρήοή one might expect borane ( B H ) to be the unstable species detected. Although no positive physical identification has been made, considerable chemical evidence of at least transitory existence is available. T o examine the compatibility of the observed spectrum with that which might be expected of borane, the authors have outlined the following argument. If one assumes that B H belongs to the point group D , there are six normal vibrational modes (see Figure 9 ) . One is inactive in the infrared, one is active and 3

3

3 h

v o 3

v

Figure 9.

Normal frequencies, X Y

3

4

b

planar molecule

nondegenerate, and two are doubly degenerate. Using equations developed by Herzberg (δ) for molecules of symmetry Z> and assuming the frequencies at 3.9, 6.40, and 8.4 microns as the normal fundamental frequencies of B H the force con stants k ks and k& for potential energy Equation 1 were obtained. (The frequency was taken as 6.40 microns because v is expected to be a doublet and may well be represented by the two peaks at 6.15 and 6.55 microns, respectively, with a center at about 6.4 microns.) 3h

3

ly

f

2

2V = h(Q

2

l2

+ Qi3 + Qit) + 2

+ Μ + «34 ) + ΜΔ12 + Δ13 + Δ ) 2

2

2

14

2

(1)

Subscripts 1, 2, 3, and 4 refer to the boron and three hydrogen atoms, respec tively; k the stretching force constant; k bending force constant i n the plane, and lf

ô>

MORREY ET AL.

165


&Δ the bending force constant out of the plane ; Q is the displacement from equilibrium of atom i from atom represents the angular displacement from equilibrium of atom i from atom ; i n the plane of symmetry and represents the angular displace ment of atom i from atom ; from equilibrium out of the plane. The resulting secular equation leads to the following equations {j

λ, - ki/m,

(2)

λ - (1 + Smy/m^ikà/niy Ρ)

(3)

λ, - f λ4 - [1 + 3/2(m„/m )] (h/m, + Zkt/mJ?)

(4)

λ λ = 3 (1 + Sm /m )(kik6/m W)

(5)

2


x

3

v

4

x

v

where A = 4π ν^ and ν is the fundamental frequency, s e c . The masses m and m refer to hydrogen and boron, respectively, and I is the distance between the boron and hydrogen atoms. Rearranged these equations give - 1

2

t

y

(6)

ki = \im

y

where

= V [B + V F - 4 C ]

fci/P

a

(7)

2

Β = my(\ + λ4)/3(1 + Smy + 2m*)

(8)

t

and C - (λ,λ477ΐ )/9(1 + 3m /wx)

(9)

2

ν

y

= 3C/(fc/P)

(10)

W P = (X m )/(1 + 3m /m,) (11) Using primed symbols to represent isotopic masses or frequencies, one can combine Equations 2 and 5 to produce equations of the Teller-Redlich product rule: 2

v

λΐ V λ

=

2

XsV XsX* k B

lt

y

My

m (l + Smy/m ') ra '(l + 3my/m ) v

β

^)

x

x

v

(1 + Smv'Mmv (1 -|- 3m /m )w ' 2

v

x

tf

( 1 4 ) 2

The measured frequencies were assigned to each of the A's and the force constants k , and & were obtained from these values and tabulated i n Table I I . Other X molecules having D symmetry are also included. 5

1 1

Δ

3

3

Table II.

Fundamental Frequencies and Force Constants of Plane Symmetrical XY3 Molecules Assuming Valence Forces v\

Molecule B»H»

B»F« B"Cl8 B»Bn

cm.

V%

vt

2560 ± · 100 1900 ± 300 1446 958 806

1560 ± · 50 1215 db 36 691 462 372

2400 db 300 1700 ± 200 888 471 279

B»Dj

β

H

VI

hi

1190=fc« 10 873 =fc 23 480 243 151

3.41 ± 0.78 3.41 ± 0.78 8.83 4.63 3.66

k /l* A

Χ 1 0 · dyne c m . -

- 1

E x p e r i m e n t a l frequencies.

1.13 ± 0.04 1.13 ± 0.04 0.87 0.42 0.29

0.25 ± 0.05 0.25 ± 0.05 0.37 0.16 0.13 - 1

U n c e r t a i n t y ie expressed i n s t a n d a r d d e v i a t i o n .

A comparison of À^'S for the various molecules shows a decrease as mass increases for the halides. The fact that k for B H is small m a y seem inconsistent. H o w ever the stretching force constant for diborane is 3.42 Χ 10 dyne c m . compared to 3.41 X 10 dyne c m . - calculated for B H . I n calculating the entropy of B H , Bauer (1) used the constants k = 3.38 X 10 , k /l = 0.10 X 10 , and k /P = 0.22 X 10 dyne c m . n

x

3

5

5

1

3

x

- 1

- 1

5

3

6

2

5

A

5


166

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B I 9 B H B I I I I I 9 I I ffi n i l J

mm

Figure 10. Infrared spectrum of B D 2

e

If one assumes the same force constants for deuterated borine, i t is possible to predict the frequencies which should show u p i f deuterated diborane were pyrolyzed u n der the same conditions that normal diborane was. These are given i n Table I I . I n order to test the prediction B D was pyrolyzed. F o r comparison purposes, a spectrum of pure B D is given i n Figure 10. Figure 11, analogous to Figures 5 and 6, shows the pyrolysis spectrum of B D . Using the same arguments as outlined i n the analysis of pyrolysis of diborane leads to results tabulated i n Table I I I . 2

2

e

e

2

Table III. Molecule BHi BDI BHI BDI

Predicted Frequency, μ 2.67 ± 0.08 3.35 ± 0.58 5.2

± 0 . 9

BDI BHI BDI B»H, B»D»

Comparative Frequencies of Isotope Molecules Possible Assignment + VA

vt vt + vi VI vt VI

BHI (8.21

11.45

±0.25)

±0.30

e

Experimental Frequency, μ 2.84 ± 0.03 3.91 ± 0 . 0 3 3.9 ± 0 . 1 5 5.35 Î6.15 I 6.55 ί 8.25 [ 8.72

vt Vi

VA

±0.03

±0.15 6.40 ± 0 . 2 0 ± 0.15 ± 0.15 8.50 ± 0.10 ± 0.15 8.40 ± 0.10 11.45 ± 0 . 0 2 5.54 7.40

Description Weak, fairly broad Weak, fairly b r o a d Probably strong, might

have

doublet structure Probably strong, might have doublet structure Possibly a doublet, strong Possibly a doublet, strong Strong, fairly wide Weak, sharp Weak, broad Weak, broad

Aside from the combination frequencies, the results agree within experimental error. Nevertheless, one wonders why the peak at 11.45 microns for B D was attenu3

WAVE

LENGTH, MICRONS

Figure 11. Infrared spectra

MORREY ET AL.

167


ated considerably in comparison with the corresponding peak of B H . I n fact, an analysis of relative absorptions had to be made to show its existence. These are summarized i n Table I V . 3

Table IV.

Effect of Time on Absorbances at 11.27, 11.45, and 11.70 (B2D6, Compensating Spectrum)

Microns


Absorbance Approximate Time, M i n . 8 15 22 Pure B D e 2

Di(11.27) 0.132 0.203 0.286 0.163

IM11.45) 0.155 0.278 0.586 0.310

Da(11.70) 0.177 0.260 0.364 0.181

D2/D1 1.17 1.37 2.05 1.91

Dz/Di 0.98 1.05 1.61 1.72

Dt/Dt 1.34 1.29 1.28 1.11

Although results are encouraging it is obvious that more refined measurements will have to be made before the assignment to B H can be considered proved. 3

Literature Cited (1) (2) (3) (4) (5) (6)

Bauer, S. H., J. Am. Chem. Soc. 78, 265 (1954). Bragg, J. K., McCarty, L. V., Norton, F. J., Ibid., 73, 2134 (1951). Burg, A. B., Schlesinger, H . I., Ibid., 55, 4009 (1933). Clarke, R. P., Pease, R. N., Ibid.,73,2132(1951). Dillard, C. R . , doctoral thesis, University of Chicago, 1949. Herzberg, G., "Molecular Spectra and Molecular Structure," Vol. II, "Infrared and Raman Spectra of Polyatomic Molecules," p. 177, Van Nostrand, New York.

RESEARCH supported by the United States Air Force through the Air Force Office of Scientific Research of the Air Research and Development Command, under contract No. AF 49(638)-28. Reproduction in whole or in part is permitted for any purpose of the United States Government.

BORAX TO BORANES

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